3.1.8 - Thermodynamics Flashcards
Define enthalpy change of formation (ΔfH)
Enthalpy change when 1 mole of a compound is formed from its constituent elements in their standard states under standard conditions
Write an equation representing the enthalpy change of formation of ethanol
Define bond dissociation enthalpy (ΔdissH)
Enthalpy change when one mole of a (covalent) bond is broken into (two) gaseous atoms (or free radicals)
Write an equation representing the bond dissociation enthalpy of chlorine
Define enthalpy change of atomisation of an element (ΔatH)
Enthalpy change when 1 mole of gaseous atoms is formed from an element in its standard state
Write an equation representing the enthalpy change of
atomisation of chlorine
Define enthalpy change of atomisation of a compound (ΔatH)
Enthalpy change when 1 mole of a compound in its standard state in converted to gaseous atoms
Write an equation representing the enthalpy change of atomisation of sodium chloride
Define first ionisation energy (Δie1H)
Enthalpy change when 1 mole of gaseous 1+ ions is formed from 1 mole of gaseous atoms
Write an equation representing the first ionisation energy of magnesium
Define second ionisation energy (Δie2H)
Enthalpy change when 1 mole of gaseous 2+ ions is formed from 1 mole of gaseous 1+ ions
Write an equation representing the second ionisation energy of magnesium
Define first electron affinity (Δea1H)
Enthalpy change when 1 mole of gaseous 1- ions is formed from 1 mole of gaseous atoms
Write an equation representing the first electron affinity of oxygen
Define second electron affinity (Δea2H)
Enthalpy change when 1 mole of gaseous 2- ions is formed from 1 mole of gaseous 1- ions
Write an equation representing the second electron affinity of oxygen
Define enthalpy change of hydration (ΔhydH)
Enthalpy change when 1 mole of aqueous ions is formed from 1 mole of gaseous ions
Write an equation representing the enthalpy change of hydration of sodium
Define enthalpy change of solution (ΔsolutionH)
Enthalpy change when 1 mole of solute is dissolved in enough solvent that no further enthalpy change occurs on further dilution
Write an equation representing the enthalpy change of solution of sodium chloride
Define lattice enthalpy of formation (ΔlatticeH)
Enthalpy change when 1 mole of solid ionic compound is formed from its gaseous ions under standard conditions
Write an equation representing the lattice enthalpy of formation of magnesium chloride
Exothermic
Define lattice enthalpy of dissociation (ΔlatticeH)
Enthalpy change when 1 mole of solid ionic compound is completely dissociated into its gaseous ions under standard conditions
Write an equation representing the lattice enthalpy of dissociation of magnesium chloride
Endothermic
What can you use to calculate lattice enthalpy since you can’t calculate it directly?
Born-Haber cycle
Calculate the lattice enthalpy of formation of sodium chloride given that the enthalpy of formation of sodium chloride is -411 kJ mol-1
Calculate the atomisation enthalpy of magnesium given that the enthalpy of formation of magnesium oxide is -548 kJ mol-1 and lattice enthalpy of formation is -3791 kJ mol-1
Calculate the lattice enthalpy of formation of aluminium oxide given that the enthalpy of formation of aluminium oxide is -1669 kJ mol-1
What does the purely ionic model of a lattice assume?
- All ions are spherical & their charge is evenly distributed around them
- Only ionic bonding
Why are experimental lattice enthalpy values different to theoretical lattice enthalpy values?
- Most ionic compounds have some partial covalent bonding
- +ve/-ve ions in lattice aren’t usually spherical
- +ve ions polarise neighbouring -ve ions to different extents
Explain why the differences between experimental and lattice enthalpies are much bigger for magnesium halides than sodium halides
- Shows bonding in magnesium halides is stronger than ionic model predicts
- ∴ bonds are strongly polarised & have quite a lot of covalent character
- Bonding in sodium halides is similar to production of ionic model = compounds closer to being purely ionic
When a solid ionic lattice dissolves in water, what 2 things occur?
- Bonds between ions break to give free ions
- Endothermic
- Bonds between ions and water are made
- Exothermic
Explain why water molecules can bond to ions
- ∵ oxygen is more electronegative than hydrogen ∴ it draws electrons towards itself = creating a dipole
- Dipole means +ve charged H atoms can form bonds with -ve ions & -ve charged O atoms can form bonds with +ve ions
Substances generally only dissolve if the energy…
released is roughly the same, or greater than the energy taken in
Souble substances tend to have ______ enthalpies of solution
exothermic
Calculate the enthalpy of solution for sodium chloride given that the lattice dissociation enthalpy is 787 kJ mol-1 and the enthalpies of hydration of sodium and chloride ions are -406 and -364 kJ mol-1 respectively.
Calculate the enthalpy of solution for silver chloride given that the lattice dissociation enthalpy is 905 kJ mol-1 and the enthalpies of hydration of silver and chloride ions are -464 and -364 kJ mol-1 respectively.
Calculate the enthalpy of solution for magnesium chloride given that the lattice formation enthalpy is -2493 kJ mol-1 and the enthalpies of hydration of magnesium and chloride ions are -1920 and -364 kJ mol-1 respectively.
What is entropy (S)?
Measure of no. of ways that particles can be arranged & no. of ways that energy can be shared out between particles
More disordered the particles are, the ____ (more +ve the value) entropy is
higher
Name and describe 2 factors that affect entropy
- Physical State
- Solid = hardly any disorder = lowest entropy
- Gas = most disordered arrangements = highest entropy
- More Particles (means more entropy)
- More particles you’ve got = more ways they + their energy can be arranged
Why do substances always tend towards disorder?
They’re more energetically stable when there’s more disorder
∴ particles will move to increase their entropy
Explain why the reaction of sodium hydrogencarbonate with hydrochloric acid is feasible even though it’s an endothermic reaction
- Feasible due to the increase in entropy as the reaction produces CO2 gas and water
- Liquids and gases are more disordered that solids and so have a higher entropy
- & more moles are being produced
- This increase in entropy overcomes the change in enthaply
- TΔS>ΔH
- ΔG<0
During a reaction, there’s an _____ ____ (ΔS) between reactants and products
Entropy change
State the equation you would use to calculate ΔS
State the units of entropy
J K-1 mol-1
Define standard entropy of a substance (S⦵)
Entropy of 1 mole of that substance under standard conditions
What is free energy change (ΔG)?
Measure used to predict whether a reaction is feasible
Free Energy Change
When is a reaction considered feasible?
If ΔG = -ve or 0
State the equation for calculating ΔG (free energy change). Include the units.
Even if ΔG shows that a reaction is theoretically feasible, it might have a … or be … that you wouldn’t notice it happening at all
Even if ΔG shows that a reaction is theoretically feasible, it might have a very high activation energy or be so slow that you wouldn’t notice it happening at all
Feasibility depends on _____
Temperature
When is ΔG always negative (i.e. reactions feasible at any temperature)? Name the 2 conditions
If reaction is exothermic (ΔH = -ve) & has +ve entropy change
When is ΔG always postive (i.e. reactions are NOT feasible at any temperature)? Name the 2 conditions
If reaction is endothermic (ΔH = +ve) & has -ve entropy change
What happens to the reaction’s feasibility if ΔH = positive & ΔS = positive?
Reaction won’t be feasible at some temps but will be at high enough temperature
What happens to the reaction’s feasibility if ΔH = negative & ΔS = negative?
Reaction will be feasible at lower temperatures but won’t be feasible at higher temperatures
Describe how you can calculate the temperature at which a reaction becomes feasible
Can find temperature when ΔG = 0
State the equation for calculating the temperature at which a reaction becomes feasible. Include the units.
Suggest why the electron affinity of chlorine is an exothermic change (2)
- Net attraction between the chlorine nucleus and the extra electron
- Energy is released (when the electron is gained)
State whether you would expect the value of the theoretical enthalpy of lattice dissociation for silver chloride to be greater than, equal to or less than that for silver bromide. Explain your answer. (3)
- Greater
- Chloride ions are smaller than bromide
- More strongly attracted to silver ions
Explain the meaning of the term perfect ionic model (1)
Ions can be regarded as point charges (or perfect spheres)
Suggest why the experimental value is greater than the theoretical value for the enthalpy of lattice dissociation for silver chloride (2)
- AgCl has covalent character
- Forces in the lattice are stronger than pure ionic attractions
Explain the interaction between water molecules and fluoride ions when the fluoride ions become hydrated (2)
- Water is polar / H on water is δ+
- (F– ions) attract water / on H δ+
Suggest why electron affinity is an endothermic process for this reaction:
Ca2+(g) + e- + S-(g) ⇒ Ca2+(g) + S2-(g)
- Negative S- ion
- Repels electron being added
Explain why the hydration enthaply of a fluoride ion is more negative than the hydration enthaply of a chloride ion (2)
- Fluoride (ions) smaller (than chloride) / have larger charge density
- So (negative charge) attracts (δ+ hydrogen on) water more strongly
By describing the nature of the attractive forces involved, explain why the value of the enthaply of hydration for a chloride ions is more negative than that for the bromide ion (3)
- Chloride (ions) are smaller (than bromide ions)
- So the force of attraction between chloride ions and water is stronger
- Chloride ions attract the δ+ on H of water / electron deficient H on water
Suggest why a value for the enthaply of solution of magnesium oxide is not found in any data books. (1)
Magnesium oxide reacts with water / forms Mg(OH)2
Suggest why using a Born-Haber cycle produces more accurate results than using mean bond enthalpies (1)
Mean bond enthalpies are from a range of compounds
